Not to be confused with Chloramine-T.
| Chloramine |
|
|
Other names
- Monochloramine[citation needed]
- Chloramide[citation needed]
- Chloroazane[citation needed]
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| Identifiers |
| CAS number |
10599-90-3 Y |
| PubChem |
25423 |
| ChemSpider |
23735 Y |
| EC number |
234-217-9 |
| KEGG |
C19359 N |
| MeSH |
chloramine |
| ChEMBL |
CHEMBL1162370 Y |
| Jmol-3D images |
Image 1 |
|
|
-
InChI=1S/ClH2N/c1-2/h2H2 Y
Key: QDHHCQZDFGDHMP-UHFFFAOYSA-N Y
|
| Properties |
| Molecular formula |
NH
2Cl |
| Molar mass |
51.476 g mol−1 |
| Appearance |
Colorless gas |
| Melting point |
−66 °C (−87 °F; 207 K) |
| Acidity (pKa) |
14 |
| Basicity (pKb) |
15 |
| Related compounds |
| Related amines |
- Dichloramine
- Nitrogen trichloride
|
| Except where noted otherwise, data are given for materials in their standard state (at 25 °C (77 °F), 100 kPa) |
| N (verify) (what is: Y/N?) |
| Infobox references |
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Chloramines are derivatives of ammonia by substitution of one, two or three hydrogen atoms with chlorine atoms: monochloramine (chloroamine, NH2Cl), dichloramine (NHCl2), and nitrogen trichloride (NCl3).[1] The term chloramine also refers to a family of organic compounds with the formulas R2NCl and RNCl2 (R is an organic group).
Monochloramine (chloroamine) is an inorganic compound with the formula NH2Cl. It is an unstable colorless liquid at its melting point of −66 °C, but it is usually handled as a dilute aqueous solution, wherein it is used as a disinfectant.
Contents
- 1 Synthesis and chemical reactions
- 1.1 Production
- 1.2 Decomposition
- 1.3 Reactions
- 2 Uses and chemical reactions
- 3 Uses in water treatment
- 3.1 Drinking water disinfection
- 3.2 Swimming pool disinfection
- 4 Removing chloramines from water
- 4.1 Dialysis
- 4.1.1 Ultraviolet light
- 4.1.2 Superchlorination
- 4.1.3 Ascorbic acid
- 4.1.4 Activated carbon
- 4.1.5 Campden tablets
- 4.1.6 Sodium thiosulfate
- 4.1.7 Other methods
- 5 Organic chloramines
- 5.1 Reduction of organic chloramines
- 6 Safety
- 7 See also
- 8 References
- 9 External links
Synthesis and chemical reactions
NH2Cl is a highly unstable compound in concentrated form. Pure NH2Cl decomposes violently above −40 °C.[2] Gaseous chloroamine at low pressures and low concentrations of chloroamine in aqueous solution are thermally slightly more stable. Chloroamine is readily soluble in water and ether, but less soluble in chloroform and carbon tetrachloride.[3]
Production
In dilute aqueous solution, chloroamine is prepared by the reaction of ammonia with sodium hypochlorite:[3]
- NH3 + OCl- ⟶ NH2Cl + HO-
This is also the first step of the Raschig hydrazine synthesis. The reaction has to be carried out in a slightly alkaline medium (pH 8.5 to 11). The acting chlorinating agent in this reaction is hypochloric acid (HOCl), which has to be generated by protonation of hypochlorite, and then reacts in a nucleophilic substitution of the hydroxo against the amino group. The reaction occurs quickest at around pH 8. At higher pH values the concentration of hypochloric acid is lower, at lower pH values ammonia is protonated to form ammonium ions NH4+, which do not react further.
The chloroamine solution can be concentrated by vacuum distillation and by passing the vapor through potassium carbonate which absorbs the water. Chloroamine can be extracted with ether.
Gaseous chloroamine can be obtained from the reaction of gaseous ammonia with chlorine gas (diluted with nitrogen gas):
- 2 NH3(g) + Cl2(g) ⟺ NH2Cl(g) + NH4Cl(s)
Pure chloroamine can be prepared by passing fluoroamine through calcium chloride:
- 2 NH2F + CaCl2 ⟶ 2 NH2Cl + CaF2
Decomposition
The covalent N-Cl bonds of chloroamines are readily hydrolyzed with release of hypochlorous acid:[4]
- RR'NCl + H2O ⟺ RR'NH + HOCl
The quantitative hydrolysis constant (K value) is used to express the bactericidal power of chloroamines, which depends on their generating hypochlorous acid in water. It is expressed by the equation below, and is generally in the range 10-4 – 10-10 (2.8 ✕ 10-10 for monochloroamine):
- K = (cRR'NH · cHOCl) / cRR'NCl
In aqueous solution, chloroamine slowly decomposes in a neutral or mildly alkaline (pH <= 11) medium:
- 3 NH2Cl ⟶ N2 + NH4Cl + 2 HCl
However, only a few percent of a 0.1 M chloroamine solution in water decompose according to the formula in several weeks. At pH values above 11, the following reaction slowly occurs:
- 3 NH2Cl + 3 OH- ⟶ NH3 + N2 + 3 Cl- + 3 H2O
In an acidic medium at pH values of around 4, chloroamine disproportionates to form dichloroamine, which in turn disproportionates at pH values below 3 to form nitrogen trichloride:
- 2 NH2Cl + H+ ⟺ NHCl2 + NH4+
- 3 NHCl2 + H+ ⟺ 2 NCl3 + NH4+
At low pH values, nitrogen trichloride dominates and between pH 3-5 dichloroamine dominates. These equilibria are disturbed by the irreversible decomposition of both compounds:
- NHCl2 + NCl3 + 2 H2O ⟶ N2 + 3 HCl + 2 HOCl
Reactions
In water chloroamine acts neutral. It is an oxidizing agent (acidic solution: E° = -1.48 V, in basic solution E° = -0.81 V):[3]
- NH2Cl + 2 H+ + 2 e- ⟶ NH4+ + Cl-
Reactions of chloroamine include radical, nucleophilic, and electrophilic substitution of chlorine, electrophilic substitution of hydrogen, and oxidative additions.
Chloroamine can, like HOCl, donate positively charged chlorine in reactions with nucleophiles:
- Nu- + NH3Cl+ ⟶ Nu–Cl + NH3
Examples of chlorination reactions include transformations to dichloroamine and nitrogen trichloride in acidic medium, as described in the decomposition section.
Chloroamine may also aminate nucleophiles (electrophilic amination):
- Nu- + NH2Cl ⟶ Nu–NH2 + Cl-
The amination of ammonia with chloroamine to form hydrazine is an example of this mechanism (Raschig process):
- NH2Cl + NH3 + NaOH ⟶ N2H4 + NaCl + H2O
Chloramine electrophilically aminates itself in neutral and alkaline media to start its decomposition:
2 NH2Cl ⟶ N2H3Cl + HCl
The chlorohydrazine (N2H3Cl) formed during self-decomposition is unstable and decomposes itself, which leads to the net decomposition reaction:
- 3 NH2Cl ⟶ N2 + NH4Cl + 2 HCl
Uses and chemical reactions
NH2Cl is a key intermediate in the traditional synthesis of hydrazine.
Monochloramine oxidizes sulfhydryls and disulfides in the same manner as HClO,[5] but only possesses 0.4% of the biocidal effect of HClO.[6]
Uses in water treatment
See also: Chloramination
Chloramine is used as a disinfectant for water because it is less aggressive than chlorine and more stable against light than hypochlorites.[3]
Drinking water disinfection
NH2Cl is commonly used in low concentrations as a secondary disinfectant in municipal water distribution systems as an alternative to chlorination. This application is increasing. Chlorine (referred to in water treatment as free chlorine) is being displaced by chloramine—to be specific monochloramine—which is much more stable and does not dissipate as rapidly as free chlorine. NH2Cl also has a very much lower, however still present, tendency than free chlorine to convert organic materials into chlorocarbons such as chloroform and carbon tetrachloride. Such compounds have been identified as carcinogens and in 1979 the United States Environmental Protection Agency began regulating their levels in U.S. drinking water.[citation needed]
Some of the unregulated byproducts may possibly pose greater health risks than the regulated chemicals.[7]
Adding chloramine to the water supply may increase exposure to lead in drinking water, especially in areas with older housing; this exposure can result in increased lead levels in the bloodstream, which may pose a significant health risk.[8]
Swimming pool disinfection
In swimming pools, chloramines are formed by the reaction of free chlorine with organic substances. Chloramines, compared to free chlorine, are both less effective as a sanitizer and, if not managed correctly, more irritating to the eyes of swimmers. Chloramines are also responsible for the reported "chlorine" smell of swimming pools.[9][10] Pool test kits designed for use by homeowners are sensitive to both free chlorine and chloramines, which can be misleading.[citation needed] There is also evidence that exposure to chloramine can contribute to respiratory problems, including asthma, among swimmers.[11] Respiratory problems related to chloramine exposure are common and prevalent among competitive swimmers.[12]
New swimming pool initially filled with tap water treated by the utility with chloramine, demonstrating how chloramine in water exhibits a greenish cast, whereas pure water is bluish (see Color of water). This greenish color may also be demonstrated by filling a white polyethylene bucket with chloraminated tap water and comparing it to chloramine-free water such as distilled water, rainwater, or a sample from a swimming pool free of chloramine.
Removing chloramines from water
Chloramines should be removed from water for dialysis, aquariums, hydroponic applications, and homebrewing beer. Chloramines can interfere with dialysis, can hurt aquatic animals, and can give homebrewed beer a medicinal taste by forming chlorophenols. In hydroponic applications, it will stunt the growth and fruit production of plants.[citation needed]
When a chemical or biological process that changes the chemistry of chloramines is used, it falls under reductive dechlorination. Other techniques use physical—not chemical—methods for removing chloramines.[citation needed]
Dialysis
Chloramine must be removed from the water prior to use in kidney dialysis machines, as it would come in contact with the bloodstream across a permeable membrane. However, since chloramine is neutralized by the digestive process, kidney dialysis patients can still safely drink chloramine-treated water.[citation needed]
Ultraviolet light
The use of UV for chlorine or chloramine removal is an established technology that has been widely accepted in pharmaceutical, beverage, and dialysis applications.[13] Ultraviolet light is also used at aquatic facilities.
Superchlorination
Chloramine can be removed from tap water by treatment with superchlorination (10 ppm or more of free chlorine, such as from a dose of sodium hypochlorite bleach or pool sanitizer) while maintaining a pH of about 7 (such as from a dose of hydrochloric acid). Hypochlorous acid from the free chlorine strips the ammonia from the chloramine, and the ammonia outgasses from the surface of the bulk water. This process takes about 24 hours for normal tap water concentrations of a few ppm of chloramine. Residual free chlorine can then be removed by exposure to bright sunlight for about 4 hours.[citation needed]
Ascorbic acid
Ascorbic acid and sodium ascorbate completely neutralizes both chlorine and chloramines but degrades in a day or two, which make it usable only for short-term applications; SFPUC determined that 1000 mg of Vitamin C (tablets purchased in a grocery store, crushed and mixed in with the bath water) remove chloramine completely in a medium-size bathtub without significantly depressing pH.[14]
Activated carbon
Activated carbon have been used for chloramine removal long before catalytic carbon became available; Standard activated carbon requires a very long contact time, which means a large volume of carbon is needed. For thorough removal, up to four times the contact time of catalytic carbon may be required.
Most dialysis units now depend on granular activated carbon (GAC) filters, two of which should be placed in series so that chloramine breakthrough can be detected after the first one, before the second one fails.[15] Additionally, sodium metabisulfite injection may be used during circumstances.[16]
Campden tablets
Home brewers use reducing agents such as sodium metabisulfite or potassium metabisulfite (both proprietary sold as campden tablets) to remove chloramine from brewing fermented beverages. However, residual sodium can cause off flavors in beer[17] so potassium metabisulfite is preferred.
Sodium thiosulfate
Sodium thiosulfate is used to dechlorinate tap water for aquariums or treat effluent from waste water treatments prior to release into rivers. The reduction reaction is analogous to the iodine reduction reaction. Treatment of tap water requires between 0.1 grams and 0.3 grams of pentahydrated (crystalline) sodium thiosulfate per 10 liters of water. Many animals are sensitive to chloramine, and it must be removed from water given to many animals in zoos.
Sodium thiosulfate has been used as treatment of calciphylaxis in hemodialysis patients with end-stage renal disease.[18]
Other methods
Chloramine, like chlorine, can be removed by boiling and aging. However, time required to remove chloramine is much longer than that of chlorine. The time required to remove half of the chloramine (half-life) from 10 gallons of water by boiling is 26.6 hours, whereas the half-life of free chlorine in boiling 10 gallons of water is only 1.8 hours.[19]
Organic chloramines
A variety of organic chloramines are known and proven useful in organic synthesis. Examples include N-chloromorpholine (ClN(CH2CH2)2O), N-chloropiperidine, and N-chloroquinuclidinium chloride.[20]
Reduction of organic chloramines
Chloramines are often an unwanted side-product of oxidation reactions of organic compounds (with amino groups) with bleach. The reduction of chloramines back into amines can be carried out through a mild hydride donor. Sodium borohydride will reduce chloramines, but this reaction is greatly sped up with acid catalysis.[citation needed]
Safety
US EPA drinking water quality standards limit chloramine concentration for public water systems to 4 parts per million (ppm) based on a running annual average of all samples in the distribution system. In order to meet EPA-regulated limits on halogenated disinfection by-products, many utilities are switching from chlorination to chloramination. While chloramination produces fewer regulated total halogenated disinfection by-products, it can produce greater concentrations of unregulated iodinated disinfection by-products and N-nitrosodimethylamine.[21][22] Both iodinated disinfection by-products and N-nitrosodimethylamine have been shown to be genotoxic.[22]
Chloroamine has been implicated as a mutagen and as a toxic agent for aquatic life, hence the US EPA proposes to prohibit its use in drinking water.[4]
See also
- Disinfection
- Disinfection by-products
- Water treatment
- Pathogen
References
- ^ Clause 2.4 Chloramines ISO 7393-2
- ^ Holleman, A. F.; Wiberg, E. "Inorganic Chemistry" Academic Press: San Diego, 2001. ISBN 0-12-352651-5.
- ^ a b c d Anton Hammerl; Thomas M. Klapötke (2005), "Nitrogen: Inorganic Chemistry", Encyclopedia of Inorganic Chemistry (2nd ed.), Wiley, pp. 55–58
- ^ a b Yasukazu Ura; Gozyo Sakata (2007), "Chloroamines", Ullmann's Encyclopedia of Industrial Chemistry (7th ed.), Wiley, p. 5
- ^ Jacangelo, J. G., V. P. Olivieri, and K. Kawata. 1987. Oxidation of sulfhydryl groups by monochloramine. Water Res. 21:1339–1344.
- ^ Morris, J. C. 1966. Future of chlorination. J. Am. Water Works Assoc. 58:1475–1482.
- ^ Stuart W. Krasner (2009-10-13). The formation and control of emerging disinfection by-products of health concern 367 (1904). Philosophical Transactions of the Royal Society. p. 4077-95. doi:10.1098/rsta.2009.0108.
- ^ Marie Lynn Miranda, et al. (February 2007). "Changes in Blood Lead Levels Associated with Use of Chloramines in Water Treatment Systems". Environmental Health Perspectives 115 (2): 221–5. doi:10.1289/ehp.9432.
- ^ Donegan, Fran J.; David Short (2011). Pools and Spas. Upper Saddle River, New Jersey: Creative Homeowner. ISBN 978-1-58011-533-9.
- ^ "Controlling Chloramines in Indoor Swimming Pools". NSW Government. Retrieved 2013-02-15.
- ^ Valérie Bougault, et al. (2009-01-01). "The Respiratory Health of Swimmers". Sports Medicine 39 (4): 295–312.
- ^ "The determinants of prevalence of health complaints among young competitive swimmers". International Archives of Occupational and Environmental Health 80 (1): 32–9. 2006-10-01.
- ^ Adelstein, Ben (2010-10-13). "Considering UV technology in water bottling". Watertechonline.com. Retrieved 2013-11-23.
- ^ "Questions Regarding Chlorine and Chloramine Removal From Water (Updated June 2013)". San Francisco Public Utilities Commission. Retrieved 2013-11-23.
- ^ Ward DM. (Oct 1996). "Chloramine removal from water used in hemodialysis". Adv Ren Replace Ther. 3 (4): 337–47. PMID 8914698.
- ^ Handbook of Dialysis, page 81
- ^ Brewing, Michael Lewis
- ^ Cicone JS, Petronis JB, Embert CD, Spector DA (June 2004). "Successful treatment of calciphylaxis with intravenous sodium thiosulfate". Am. J. Kidney Dis. 43 (6): 1104–8. doi:10.1053/j.ajkd.2004.03.018. PMID 15168392.
- ^ "Experiments in Removing Chlorine and Chloramine From Brewing Water" (PDF). 1998-11-03. Retrieved 2013-11-23.
- ^ Lindsay Smith, J. R.; McKeer, L. C.; Taylor, J. M. "4-Chlorination of Electron-Rich Benzenoid Compounds: 2,4-Dichloromethoxybenzene" Organic Syntheses, CollectedVolume 8, p.167 (1993)..http://www.orgsyn.org/orgsyn/pdfs/CV8P0167.pdf describes several N-chloramines
- ^ Krasner, Stuart W.; Weinberg, Howard S.; Richardson, Susan D.; Pastor, Salvador J.; Chinn, Russell; Sclimenti, Michael J.; Onstad, Gretchen D.; Thruston, Alfred D. (2006). "Occurrence of a New Generation of Disinfection Byproducts". Environmental Science & Technology 40 (23): 7175–85. doi:10.1021/es060353j.
- ^ a b Richardson, Susan D.; Plewa, Michael J.; Wagner, Elizabeth D.; Schoeny, Rita; DeMarini, David M. (2007). "Occurrence, genotoxicity, and carcinogenicity of regulated and emerging disinfection by-products in drinking water: A review and roadmap for research". Mutation Research/Reviews in Mutation Research 636 (1–3): 178–242. doi:10.1016/j.mrrev.2007.09.001. PMID 17980649.
External links
- "Chlorinated drinking water", IARC Monograph (1991)
- EPA Maximum Contaminant Levels
- WebBook page for NH2Cl
- Chlorine and chloramines in the freshwater aquarium
